2.6.2: 2.5.2 Non-Aqueous Acid/Base Systems

We ordinarily think of Brønsted-Lowry acid-base reactions as taking place in aqueous solutions, but this need not always be the case. A more general view encompasses a variety of acid-base solvent systems, of which the water system is only one (Table \(\PageIndex{1}\)). Each of these has as its basis an amphiprotic solvent (one capable of undergoing autoprotolysis), in parallel with the familiar case of water.

The limiting acid in a given solvent is the solvonium ion, such as H₃O⁺ (hydronium) ion in water. An acid which has more of a tendency to donate a hydrogen ion than the limiting acid will be a strong acid in the solvent considered, and will exist mostly or entirely in its dissociated form. Likewise, the limiting base in a given solvent is the solvate ion, such as OH⁻ (hydroxide) ion, in water. A base which has more affinity for protons than the limiting base cannot exist in solution, as it will react with the solvent.

Figure } \PageIndex{3}\): Use of non-aqueous solvents allows the study of strong acids that are hindered by the "leveling effect" of the solvent.

The ammonia system is one of the most common non-aqueous system in Chemistry. Liquid ammonia boils at –33° C, and can conveniently be maintained as a liquid by cooling with dry ice (–77° C). It is a good solvent for substances that also dissolve in water, such as ionic salts and organic compounds since it is capable of forming hydrogen bonds. Bases can exist in solution in liquid ammonia which cannot exist in aqueous solution: this is the case for any base which is stronger than the hydroxide ion, but weaker than the amide ion \(NH_2^-\). The limiting acid in liquid ammonia is the ammonium ion, which has a pK_a value in water of 9.25. The limiting base is the amide ion, NH₂⁻. This is a stronger base than the hydroxide ion and so cannot exist in aqueous solution. The pK_a value of ammonia is estimated to be approximately 33. Any acid which is a stronger acid than the ammonium ion will be a strong acid in liquid ammonia. This is the case for acetic acid, which is completely dissociated in liquid ammonia solution. The addition of pure acetic acid and the addition of ammonium acetate have exactly the same effect on a liquid ammonia solution: the increase in its acidity: in practice, the latter is preferred for safety reasons.

Note

As with \(OH^-\) and \(H_3O^+\) in water, the strongest acid and base in \(NH_3\) is dictated by the corresponding autoprotolysis reaction of the solvent:

\[2 NH_3 \rightleftharpoons NH_4^+ + NH_2^– \nonumber\]

One use of non-aqueous acid-base systems is to examine the relative strengths of the strong acids and bases, whose strengths are "leveled" by the fact that they are all totally converted into H₃O⁺ or OH^– ions in water. By studying them in appropriate non-aqueous solvents which are poorer acceptors or donors of protons, their relative strengths can be determined. Many familiar substances can serve as the basis of protonic solvent systems (Table \(\PageIndex{1}\)).

\(\textbf{Table } \PageIndex{1}\): Popular Solvent systems

solvent	autoprotolysis reaction	pKa
water	\(2 H_2O \rightleftharpoons H_3O^+ + OH^–\)	14
ammonia	\(2 NH_3 \rightleftharpoons NH_4^+ + NH_2^–\)	33
acetic acid	\(2 CH_3COOH \rightleftharpoons CH_3COOH_2^+ + CH_3COO^–\)	13
ethanol	\(2 C_2H_5OH \rightleftharpoons C_2H_5OH_2^+ + C_2H_5O^–\)	19
hydrogen peroxide	\(2 HO-OH \rightleftharpoons HO-OH_2^+ + HO-O^–\)	13
hydrofluoric acid	\(2 HF \rightleftharpoons H^2F^+ + F^–\)	10
sulfuric acid	\(2 H_2SO_4 \rightleftharpoons H3SO_4^+ + HSO_4^–\)	3.5

The extreme case is a superacid, a medium in which the hydrogen ion is only very weakly solvated. The classic example is a mixture of antimony pentafluoride and liquid hydrogen fluoride:

\[SbF_5 + HF \rightleftharpoons H^+ + SbF_6^−\]

The limiting base, the hexfluoroantimonate anion \(SbF_6^−\), is so weakly attracted to the hydrogen ion that virtually any other base will bind more strongly: hence, this mixture can be used to protonate organic molecules which would not be considered bases in other solvents. It should noted that pH is undefined in aprotic solvents, which assumes presence of hydronium ions. In other solvents, the concentration of the respective solvonium/solvate ions should be used (e.g., \([NH_4^+]\) and \([NH_2^–]\) in \(NH_{3(l)}\).

Contributors and Attributions

• Stephen Lower, Professor Emeritus (Simon Fraser U.) Chem1 Virtual Textbook

• Wikipedia